Isotopes - Chemistry LibreTexts Isotopes Contributors Atoms that have the same atomic number (number of protons), but different mass numbers (number of protons and neutrons) are called isotopes. There are naturally occurring isotopes and isotopes that are artificially produced. Isotopes are separated through mass spectrometry; MS traces show the relative abundance of isotopes vs. mass number (mass : charge ratio). Introduction As mentioned before, isotopes are atoms that have the same atomic number, but different mass numbers. Isotopes are denoted the same way as nuclides , but they are often symbolized only with the mass numbers because isotopes of the same element have the the same atomic number. Carbon, for example, has two naturally occurring isotopes, \(^{12}_6C\) and \(^{13}_6C\). Because both of these isotopes have 6 protons, they are often written as \(^{12}C\) and \(^{13}C\). \(^{12}C\) has 6 neutrons, and \(^{13}C\) has 7 neutrons. Of all the elements on the periodic table, only 21 are pure elements.

Pure, or monotopic, elements are those elements with only one naturally occurring nuclide. The following lists the 21 pure elements: \(^{27}_{13}Al\) Isotopes of the other elements either occur naturally or are artificially produced. Natural and Artificial Isotopes Most elements have naturally occurring isotopes. Percent natural abundances indicate which isotopes of any given element are predominant (occur in greater abundance) and which only occur in trace amounts. Mercury, for example, has seven naturally occurring isotopes: \(^{196}Hg\), \(^{198}Hg\), \(^{199}Hg\), \(^{200}Hg\), \(^{201}Hg\), \(^{202}Hg\), \(^{204}Hg\); these have the percent natural abundances of 0.146%, 10.02%, 16.84%, 23.13%, 13.22%, 29.80%, and 6.85%, respectively.

It is clear that \(^{202}Hg\) occurs with greatest abundance, and \(^{200}Hg\) is the next most abundant, but the other isotopes only occur in small traces. Note: The sum of the percent natural abundances of all the isotopes of any given element must total 100%. There are 20 elements with only artificially produced isotopes. The majority of these are heavier elements; the lightest elements with artificial isotopes are \(^{43}Tc\) and \(^{61}Pm\). The other elements that only have artificial isotopes are those with atomic numbers of 84-88 and 89-103, otherwise known as the actinoids, but excluding \(^{90}Th\) and \(^{92}U\). Some naturally occurring and artificially produced isotopes are radioactive. The nucleus of a radioactive isotope is unstable; radioactive isotopes spontaneously decay, emitting alpha, beta, and gamma rays until they reach a stability, usually in the state of a different element. Bismuth (\(^{209}_{83}Bi\)) has the highest atomic and mass number of all the stable nuclides.

All nuclides with atomic number and mass number greater than 83 and 209, respectively, are radioactive. However, there are some lighter nuclides that are radioactive. For example, hydrogen has two naturally occurring stable isotopes, \(^{1}H\) and \(^{2}H\) (deuterium), and a third naturally occurring radioactive isotope, \(^{3}H\) (tritium). Radioisotope Dating The presence of certain radioisotopes in an object can be used to determine its age. Carbon dating is based on the fact that living plants absorb stable \(^{12}C\), \(^{13}C\) and radioactive \(^{14}C\) from the atmosphere, and animals absorb them from the plants. An organism no longer absorbs carbon after it dies, its age can be determined by measuring the ratio of \(^{13}C\) to \(^{14}C\) in the sample and extrapolating based on its decay rate. Art forgeries are often detected by similar means. \(^{137}Cs\) and \(^{90}Sr\) do not occur naturally and are only present in the atmosphere today because of nuclear weapons.

Any object created before July 1945, then, would have neither of these elements, so finding them through mass spectrometry or other means would indicate that it was created later. Isotopic Masses, Percent Natural Abundance, and Weighted-Average Atomic Mass Because most elements occur as isotopes and different isotopes have different masses, the atomic mass of an element is the average of the isotopic masses, weighted according to their naturally occurring abundances; this is the mass of each element recorded on the periodic table, also known as the relative atomic mass (Ar). Treating isotopic masses in weighted averages gives greater importance to the isotope with greatest percent natural abundance. Below is a general equation to calculate the atomic mass of an element based on percent natural abundance and isotopic masses: * fractional abundance is the percent abundance divided by 100% Bromine has two naturally occurring isotopes: bromine-79 has a mass of 78.9183 u and an abundance of 50.69%, and bromine-81 has a mass of 80.92 u and an abundance of 49.31%.

The equation above can be used to solve for the relative atomic mass of bromine: atomic mass of Br = (0.5069 x 78.9183 u) + (0.4931 x 80.92 u) = 79.91 u This is the relative atomic number of bromine that is listed on the periodic table. Comparing their isotopic masses of any given element to the relative atomic mass of the element reveals that the Ar is very close to the isotope that occurs most frequently. Thus, the isotope whose isotopic mass is closest to the atomic mass of the element is the isotope that occurs in the greatest abundance. Mass Spectrometry Mass spectrometry is a technique that can be used to distinguish between isotopes of a given element. A mass spectrometer separates each isotope by mass number. Each isotope is characterized by a peak (of given intensity) according to its relative abundance. The most intense peak corresponds to the isotope that occurs in the largest relative natural abundance, and vice versa. Refer to Mass Spectrometry: Isotope Effects . Example: The mass spectrum of strontium has four different peaks, varying in intensity.

The four peaks indicate that there are four isotopes of strontium. The four isotopes of strontium have isotopic mass numbers of 84, 86, 87, and 88, and relative abundances of 0.56%, 9.86%, 7.00%, and 82.58%, respectively. The intensity of the peak corresponds to the abundance. \(^{84}Sr\) has the smallest peak, which corresponds to its relative abundance of 0.56%, whereas \(^{88}Sr\) has the largest peak, which corresponds to its relative abundance of 82.58%. This indicates that \(^{88}Sr\) is the isotope that occurs in highest amounts. References Petrucci, Ralph H., William S. Harwood, F. Geoffrey Herring, and Jeffry D. Madura. General Chemistry: Principles and Modern Application. Ninth ed. New Jersey: Pearson Prentice Hall, 2007. Housecraft, Catherine E. and Alan G. Sharpe. Inorganic Chemistry. Third ed. England: Pearson Prentice Hall, 2008.

Hoefs, Jochen. Stable Isotope Geochemistry. Sixth ed. Germany: Springer, 2009.

Problems 1) Find the number of protons and neutrons in the following isotopes: a) \(^{20}Ne\), \(^{21}Ne\), \(^{22}Ne\); b) \(^{84}Sr\), \(^{86}Sr\), \(^{87}Sr\), \(^{88}Sr\); c) \(^{102}Pd\), \(^{104}Pd\), \(^{105}Pd\), \(^{106}Pd\), \(^{108}Pd\), \(^{110}Pd\) (Answer: a) \(^{20}Ne\) has 10 p and 10 n; \(^{21}Ne\) has 10 p and 11 n; \(^{22}Ne\) has 10 p and 12 n; b) \(^{84}Sr\) has 38 p and 46 n, \(^{86}Sr\) has 38 p and 48 n, \(^{87}Sr\) has 38 p and 49 n,

 \(^{88}Sr\) has 38 p and 50 n; c) \(^{102}Pd\) has 46 p and 56 n, \(^{104}Pd\) has 46 p and 58 n, \(^{105}Pd\) has 46 p and 59 n, \(^{106}Pd\) has 46 p and 60 n, \(^{108}Pd\) has 46 p and 62 n, \(^{110}Pd\) has 46 p and 64 n) 2) An isotope of lead has an isotopic mass of 208.

Give the appropriate notation for the isotope. (Answer: \(^{208}Pb\)) 3) Thallium has two naturally occurring istotopes, \(^{203}Tl\) and \(^{205}Tl\). Tl-205 has an abundance of 70.48%. What is the percent natural abundance of Tl-203? (Answer: 29.52%) 4) Chlorine has two naturally occuring isotopes, \(^{35}Cl\) and \(^{37}Cl\), with masses of 34.97 u and 36.97 u, respectively. Which of these two isotopes occur in greater abundance? (Answer: \(^{35}Cl\)) 5. Potassium has three naturally occurring isotopes: \(^{39}K\), \(^{40}K\), and \(^{41}K\), with masses (and abundances) of 38.96 u (93.3%), 39.96 u (0.012%), and 40.96 u (6.7%), respectively.

Calculate the relative atomic mass of K. (Answer: 39.10 u) Review Questions What do the terms 'monotopic' and 'radioactive' mean? On the mass spectrum of any given atom, what does the number and intensity of the peak(s) tell you? What is the relationship between isotopic masses, percent natural abundance, and weighted-average atomic mass? Contributors There are no recommended articles. Article type Lightest naturally occurring element - Science Mathematics Lightest naturally occurring element HOME > > Lightest naturally occurring element Lightest naturally occurring element [From: ] [author: ] [Date: 12-07-08] [Hit: ] Which of the following is the lightest naturally occurring element found on Earth? A Hydrogen What Is the Heaviest Element? What Is the Heaviest Element? What Is the Heaviest Element? Why It's Hard to Identify the Element with Highest Density Osmium is generally considered the heaviest element, but sometimes iridium surpasses it. Periodictableru Question: What Is the Heaviest Element? Are you wondering which element is the heaviest? Here's an answer to the question.

Answer: There are two ways of answering this question, depending on how you define 'heaviest'. Heaviest Element in Terms of Atomic Weight The heaviest element in terms of heaviest per a given number of atoms is the element with the highest atomic weight. This is the element with the largest number of protons, which is presently element 118, oganesson or ununoctium . When a heavier element is discovered (e.g., element 120), then that will become the new heaviest element. Ununoctium is the heaviest element, but it is man-made. The heaviest naturally-occurring element is uranium (atomic number 92, atomic weight 238.0289). Heaviest Element in Terms of Density Another way to look at heaviness is in terms of density, which is mass per unit volume. Either of two elements can be considered the element with the highest density : osmium and iridium . continue reading below our video 10 Best Universities in the United States The density of the element depends on many factors, so there isn't a single number for density that would allow us to identify one element or the other as the most dense.

Each of these elements weigh approximately twice as much as lead. The calculated density of osmium is 22.61 g/cm3 and the calculated density of iridium is 22.65 g/cm3, though the density of iridium has not been experimentally measured to exceed that of osmium. Why Osmium and Iridium Are So Heavy Even though there are many elements with higher atomic weight values, osmium and iridium are the heaviest. This is because their atoms pack together more tightly in solid form. The reason for this is that their f electron orbitals are compacted when n=5 and n=6. The orbitals feel the attraction of the positive-charged nucleus because of this, so the atom size contracts. Relativistic effects also play a role. The electrons in these orbitals go around the atomic nucleus so fast their apparent mass increases. When this happens, the s orbital shrinks. Is Uranium the Heaviest Natural Element? – Starts With A Bang Last night, I was watching the Daily Show, and they had Tom Zoellner on, talking about his new book: Uranium: War, Energy, and the Rock that shaped the world .

There are certainly a lot of interesting things about Uranium in culture, particularly in terms of energy (hey, we can use this thing to power the world) and in terms of war (the nuclear arms race). If you look at what we’ve done with all the enriched Uranium (U-235, the fissionable type), the US and the Soviets, from the 1940s to the 1980s, basically took nearly all of it and stockpiled it. This is the only element on Earth we’ve ever done that to, and that makes it incredibly interesting from a cultural point of view. But Tom Zoellner said something interesting during the interview about Uranium, that I’d like for you to think about with me: [Uranium] is one of these things… it’s the heaviest naturally occurring mineral in the periodic table. I mean, the table goes up higher, but this is where it ends. This stuff seems to hit some invisible wall of nature where it has this really big nucleus, this nucleus of 92 protons, and the center can’t hold. Now, this is mostly true. It is the heaviest naturally occurring element on Earth.

We don’t find any elements with 93 protons (Neptunium), 94 protons (Plutonium), or higher numbers except the ones we’ve made in laboratories. Here’s a version of the periodic table of naturally occurring elements for you: But — here’s a great question for you — why? Why is Uranium the heaviest element on Earth? Well, there are two reasons: one is obvious, and one is very subtle. The obvious reason is that the elements shown above in light grey, 43-Tc, 61-Pm, and everything above and including 93-Np are unstable. They get created in stars and supernovae, just like all the other elements, but they radioactively decay into the other, stable elements (1-42, 44-60, and 62-92). But hang on a second. Isn’t Uranium, element 92, also unstable? Its most stable isotope, Uranium 238 (92 protons + 146 neutrons), has a half-life of about 4.5 billion years , or the age of the Earth.

Well, that explains why there’s still so much Uranium around: it’s only had enough time for about 50% of the atoms to decay. The fissionable type of Uranium, Uranium 235 (3 fewer neutrons), is less stable, with a half-life of 700 million years , is still around, too. Much more of it — about 98.8% — is gone by now. However, there’s still enough left that there’s plenty on Earth. But every element heavier than lead (element number 82) is unstable, and will eventually decay! Of all the unstable elements that aren’t on Earth, though, I’m curious about why there isn’t any Plutonium? Plutonium-244 , with 94 protons and 150 neutrons, has a half-life of 83 million years, which means that there should still be small amounts of it left today. Plutonium is more stable than Protactinium and Actinium, which are found in Nature. Only one study, whose results are not universally accepted, claims to have detected natural Plutonium .

So, now for the more subtle point: why isn’t there any Plutonium-244 on Earth? The only way you produce elements heavier than Iron-56 is in Supernovae: when massive, dying stars collapse and explode. These heavy elements then get spread out over a large area and recollapse, forming new stars and planets. That’s where all of the elements on Earth came from. But the supernova that gave rise to us must not have been powerful enough to make huge amounts of Plutonium! Since there are supernovae that are up to 100 times more powerful than normal ones , there are explosions out there in the Universe that will litter it with Plutonium. We saw one in 2006 that probably did just that: Well, the planets that get created from the dust of this explosion will not see Uranium (element 92) as their “wall of nature,” they’ll see Plutonium (element 94) as the heaviest, or possibly even Curium-247 (element 96)! There are people searching for this as well.

And, on a lighter note, there will surely be planets out there where the heaviest element is lighter than Uranium, possibly even lighter than lead! So yes, Uranium is the heaviest natural element that we find on Earth, but there are certainly heavier ones out there! Element 118, Heaviest Ever, Reported for 1,000th of a Second - The New York Times The New York Times Science |Element 118, Heaviest Ever, Reported for 1,000th of a Second Search Continue reading the main story A team of Russian and American scientists said yesterday that it had created the heaviest element ever seen in a laboratory, a dab of matter that lasted for less than one-thousandth of a second but would add an entry at the farthest reaches of the periodic table and suggest that strange new elements may lie beyond. By convention, the substance remains the Baby Doe of elements until its existence is confirmed at other laboratories. For now, the new substance will be principally known as element 118 for the number of protons in its nucleus, more than in any other element occurring naturally or produced in the laboratory.

Hydrogen, the lightest element, has one proton in its nucleus, and uranium, the heaviest naturally occurring element, has 92. Element 118 would fit comfortably just below radon in a column of the periodic table containing what are called noble gases for their inert chemical properties. The results were met with praise but also caution from other scientists in the field, particularly given the fraught history of element 118. Another California lab, the Lawrence Berkeley National Laboratory, announced that it discovered the element in 1999 but retracted the claim two years later after an investigation found that one of its researchers, Dr. Victor Ninov, had fabricated data. Dr. Ninov was later fired. Continue reading the main story Dr. Ken Moody, the lead American researcher on the work, said everything had been done to guard against fabrication, with independent analyses being carried out in Russia and the United States. And the group’s paper on the putative find has been accepted at a prestigious journal, Physical Review C, after other scientists reviewed the work, said Dr. Jonathan Lenaghan, an editor at the journal.

Advertisement Continue reading the main story But it was less the fear of fraud than ordinary scientific caution that caused some scientists to reserve judgment on the discovery. The Russian lab and its collaborators have now announced the discovery of five elements — 113, 114, 115, 116 and now 118 — none of which have been confirmed by other scientists. “One has to be extremely careful with those enthusiastic announcements,” said Dr. Witold Nazarewicz, a nuclear theorist at the University of Tennessee and the Oak Ridge National Laboratory. “This is not because one is doing something wrong,” Dr. Nazarewicz said. “It’s because these are very difficult measurements. They are playing on the edges of statistics.” Photo Calcium, with 20 protons, being accelerated into Californium, with 98 protons. Credit Sabrina Fletcher and Thomas Tegge/Lawrence Livermore National Laboratory The team that created the element, made up of scientists at the Lawrence Livermore National Laboratory in California and the Joint Institute for Nuclear Research in Dubna, Russia, said they had produced three atoms of the new element in six months of smashing lighter elements together and trying to make them stick.

The scientists said their results also gave hope that they were approaching a long-predicted “island of stability” of even heavier elements, with longer lives and possibly strange new chemical properties. “This considerably expands the borders of the existing material world,” Dr. Yuri Oganessian said in an e-mail message. He is the scientific director of the Flerov Laboratory of Nuclear Reactions at the Dubna institute, where the experiments were carried out in a cyclotron, the circular particle accelerator. A Livermore scientist, Dr. Nancy J. Stoyer, said the team had calculated that there was less than one chance in 100,000 that the results were a statistical fluke. “We’re very confident,” Dr. Stoyer said. Other experimental scientists said nothing was obviously amiss with the work. Please verify you're not a robot by clicking the box. Invalid email address. Please re-enter. You must select a newsletter to subscribe to. Sign Up “I think the evidence they have is convincing,” said Dr. C. Konrad Gelbke, director of the National Superconducting Cyclotron Laboratory at Michigan State University.

“It looks pretty good.” Advertisement Continue reading the main story The experiments were performed when scientists at the Russian laboratory used the cyclotron to bash atoms of calcium, with 20 protons, into a target of Californium, with 98 protons, like little clumps of putty that they hoped would stick together, said Dr. Dawn A. Shaughnessy, another Livermore scientist who worked on the experiment. In extremely rare instances, they did stick. In 10 billion billion bombardments, detectors found that the two sets of protons glommed together to produce element 118. Photo The new element traveling through the accelerator to the detector. Credit Sabrina Fletcher and Thomas Tegge/Lawrence Livermore National Laboratory An element’s weight is determined by the total number of protons, which have a positive electrical charge, and neutrons, which are neutral, in its nucleus. By that measure, too, the new element is the heaviest ever created. Dr. Gelbke said that there was one good reason that the Russian laboratory might be ahead of its competitors elsewhere in the world.

Scientists at that lab, he said, are skilled in handling Californium, which is very radioactive and dangerous to the uninitiated. “I wouldn’t want to do that myself,” Dr. Gelbke said, chucking. “It’s a fairly nasty substance for most people to handle.” If the results are confirmed, they would represent one more step toward the “island of stability” that theorists have predicted in even heavier regions of the periodic table. Nuclei have shell-like structures, and the most stable atoms contain so-called “magic numbers” of protons and neutrons that produce closed, or complete, shells. The numbers 2, 8, 20, 28, 50 and 82 are magic for both protons and neutrons. The highest known magic number for neutrons alone is 126, meaning that common lead, with 82 protons and 126 neutrons, is the heaviest known “doubly magic,” or extremely stable, isotope in the periodic table. But the theorists have predicted that there is another closed shell out beyond all elements discovered so far, including the latest one.

Advertisement Continue reading the main story “It’s rather like Plum Island at the end of Long Island,” Dr. Martin Blume, the overall editor of Physical Review, said. “You go there, there’s a gap, and then there’s Plum Island.” There is general agreement that the next neutron magic number is 184. But that is still out of reach of current experiments. The next proton magic number is a matter of disagreement. “That is, I think, the basis for looking in this region,” Dr. Blume said. “Have you reached the island of stability?” A version of this article appears in print on , on Page F1 of the New York edition with the headline: Element 118, Heaviest Ever, Reported for 1,000th of a Second. Order Reprints | Today's Paper | Subscribe How elements are formed | Sciencelearn Hub How elements are formed Image: Neon sign Our world is made of elements and combinations of elements called compounds. An element is a pure substance made of atoms that are all of the same type. At present, 116 elements are known, and only about 90 of these occur naturally.

Elements and the ‘Big Bang’ theory During the formation of the universe some 14 billion years ago in the so-called ‘ Big Bang ’, only the lightest elements were formed – hydrogen and helium along with trace amounts of lithium and beryllium. As the cloud of cosmic dust and gases from the Big Bang cooled, stars formed, and these then grouped together to form galaxies. The other 86 elements found in nature were created in nuclear reactions in these stars and in huge stellar explosions known as supernovae. Elements and our Sun Image: The Sun For most of their lives, stars fuse elemental hydrogen into helium in their cores. Two atoms of hydrogen are combined in a series of steps to create helium-4. These reactions account for 85% of the Sun’s energy. The remaining 15% comes from reactions that produce the elements beryllium and lithium. The energy from these nuclear reactions is emitted in various forms of radiation such as ultraviolet light, X-rays , visible light, infrared rays, microwaves and radio waves . In addition, energised particles such as neutrinos and protons are released, and it is these that make up the solar wind .

Earth is in the path of this energy stream, which warms the planet , drives weather and provides energy for life. The Earth’s atmosphere is able to screen out most of the harmful radiation, and the Earth’s magnetic field can deflect the harmful effects of the solar wind. Dying stars When a star ’s core runs out of hydrogen, the star begins to die out. The dying star expands into a red giant , and this now begins to manufacture carbon atoms by fusing helium atoms. More massive stars begin a further series of nuclear burning or reaction stages. The elements formed in these stages range from oxygen through to iron . During a supernova , the star releases very large amounts of energy as well as neutrons, which allows elements heavier than iron, such as uranium and gold , to be produced. In the supernova explosion, all of these elements are expelled out into space. Video: What is the Big Bang theory? Our world is literally made up of elements formed deep within the cores of stars now long dead.

As Britain’s Astronomer Royal Sir Martin Rees said, “We are literally the ashes of long dead stars.” When you buy a party balloon that floats in air, it is filled with helium gas – most of which was created when the universe was only 3 minutes old! Examples of element making (nucleogenesis) in helium burning reactions: 3 helium atoms fusing to give a carbon atom : 3 @ 4He → 12C carbon atom + helium atom fusing to give an oxygen atom: 12C + 4He → 16O oxygen atom + helium atom fusing to give a neon atom: 16O + 4He → 20Ne neon atom + helium atom fusing to give a magnesium atom: 20Ne + 4He → 24Mg Man-made elements Only 90 of the 116 known elements occur naturally, so where have the other 26 come from?

The answer is to be found in the development of nuclear power plants and machines known as particle accelerators: Scientists discovered that, by allowing fast neutrons to collide with the common isotope of uranium known as U-238 in a nuclear reactor , the ‘new’ element plutonium was made. By smashing atoms together in machines known as particle accelerators, it was discovered that new elements could be made. For example, bombarding atoms of the element curium with atoms of neon made element 106 – seaborgium. Nature of Science One of the habits of scientists is open-mindedness. Scientists need to be receptive to new ideas and suggestions. As new evidence is discovered, new ways of interpreting and understanding it may have to be considered. Metadata NOVA - Official Website | The Star In You NOVA "Our planet, our society, and we ourselves are built of star stuff." —Carl Sagan, Cosmos Here's an amazing fact for your next cocktail party: Every single atom in your body—the calcium in your bones, the carbon in your genes, the iron in your blood, the gold in your filling—was created in a star billions of years ago. All except atoms of hydrogen and one or two of the next lightest elements.

They were formed even earlier, shortly after the Big Bang began 13.7 billion years ago. As hard as it might be to believe, every atom in your body, astrophysicists say, originated billions of years ago in a star or in the explosive aftermath of the Big Bang. Here, a close-up of Polaris, the North Star. Enlarge Photo credit: NASA, ESA, G. Bacon (STScI) It's true, according to astrophysicists. You and everything around you, every single natural and man-made thing you can see, every rock, tree, butterfly, and building, comprises atoms that originally arose during the Big Bang or, for all but the lightest two or three elements, from millions of burning and exploding stars far back in the history of the universe. You live because stars died; it's that simple. How is this so? How can you possibly be a walking galaxy of fossil stardust? Well, the story is not a new one, but it bears retelling, if only because its working out was one of the finest achievements of 20th-century astrophysics—and because it's so astonishing.

The start of it all The story begins at the beginning, as in the Big Bang. That is when, astrophysicists say, all the hydrogen in the universe came into being. Initially it was just protons, and then, as the young universe expanded and cooled, these became bound to electrons, forming hydrogen atoms. The very hydrogen atoms in the H2O that makes up over half your body were born then. They didn't come from your parents; they came from the early universe. Did you have any idea you have atoms in your body that are over 13 billion years old? If you could separate one hydrogen atom from one molecule of water in your body, shrink down to its atomically tiny size like the scientists in Fantastic Voyage , then reverse time and follow it back to through its unimaginable lifetime, you would find yourself in the immediate aftermath of the Big Bang. That very hydrogen atom, an atom now inside you as you read this, has remained unchanged since the beginning of time. Over 13 billion years since the Big Bang, hydrogen and helium still make up most of the visible matter in the universe. Nearly 10,000 galaxies appear in this Hubble Ultra Deep Field image.

Enlarge Photo credit: NASA, ESA, and N. Pirzkal (STScI/ESA) The Big Bang also churned out helium, the next lightest element. You don't have any helium in you, unless you just sucked the gas out of a birthday balloon. But helium is the second most common element after hydrogen. Together they make up more than 98 percent of the matter in the universe. (Luminous matter, that is; dark matter is a whole other story.) A smattering of lithium (element 3) and one or two other of the lightest elements also formed in the Bang, but these were negligible. Everything else, every other chemical element, including carbon, oxygen, nitrogen, and all the other elements essential for your life, is thought to have been fabricated in stars. How? Well, the story is either simple or horrendously complex depending on whether you're a science writer or a scientist. Here's the simple story: Table for 118 First, what are we talking about when we talk about an element? A chemical element is a substance that cannot be broken down or changed into another substance using chemical means. It can be changed using nuclear means, which is what happens inside stars.

Every second, the sun converts about 500 million tons of hydrogen into helium. As we learn in high school chemistry—and can remind ourselves with a quick glance at the Periodic Table —hydrogen, the lightest element, has one proton in its nucleus and thus is given the atomic number 1. Helium has two protons and so is number 2, and so on all the way up to uranium, which, with 92 protons in its nucleus, is the heaviest of the "naturally occurring" elements. Remarkably, all life on Earth, all everything we see around us, consists of various combinations of those 92 elements. There are still heavier elements, ranging from neptunium (93) all the way up to the unofficially named ununoctium (118), though with the exception of trace amounts of neptunium and plutonium (94), these are not found naturally on Earth. An artist's impression of how the very early universe—less than one billion years old—might have looked during an intense period of hydrogen conversion into myriad stars Enlarge Photo credit: Science: NASA and K. Lanzetta (SUNY).

Art: Adolf Schaller for STScI. Stars are born How did—and do, for the process continues today—all the chemical elements first come into existence? Several hundred million years after the Big Bang, about 13 billion years ago, the hydrogen and helium in the early universe began coalescing into gas clouds, which, in turn, collapsed into the first stars. Gravity, that not-to-be-denied force, caused these newborn stars to contract, heating their cores to temperatures high enough to ignite their hydrogen and trigger its fusion into helium. This is the first link in a chain of thermonuclear reactions that, depending on the size of the star and its fate, bring about the genesis of all the other chemical elements up to about californium, element 98. (Heavier elements than that are produced only in particle accelerators, physicists believe.) Imagine starting out in your kitchen with just a single natural ingredient and, after baking it in your oven, winding up with all other possible natural ingredients. This is what the universe has done with hydrogen. The burning of H to He is what our star, the sun, does for a living.

In the searing heat of its core—about 27 million °F—the reaction of four hydrogen nuclei fusing to become one helium nucleus happens over and over and over again, ad infinitum. Every second, the sun converts about 500 million tons of hydrogen into helium. (And for every helium atom formed, roughly a trillion photons are emitted from the sun's surface. This is why we wear sunglasses.) Cooking elements Our star enables us to live, but at this stage in its own life, it doesn't give us any elements heavier than helium. It's not massive enough. With stars more massive than ours, and up to about eight times its mass * , gravity is forcible enough to compress the core sufficiently to trigger nuclear reactions that produce heavier elements, starting with carbon (element 6) and oxygen (8). In such cores, the heat is high enough, about 180 million °F, to force three helium nuclei to fuse into a carbon nucleus, or four helium nuclei into an oxygen nucleus, millions of times over. This will happen in the sun when it becomes a red giant in five billion years.

In its fiery core, our star, the sun, produces only a single chemical element—helium—over and over again. Enlarge Photo credit: NASA In very massive stars, those of more than eight solar masses, the force of gravity drives the temperature in the core up so outlandishly high that it triggers thermonuclear reactions that create elements all the way up to iron (26). At 1,080 million °F, carbon fuses into neon; at 2,700 million °F, oxygen fuses into silicon; and at 7,200 million °F, silicon fuses into iron. Iron, alas, marks a major turning point when it comes to fusing ever-heavier elements inside stars. All the way up to iron, every time a new fusion reaction occurs, some heat is released. With iron, no other rearrangement of nuclei can generate any more energy. But stars do form elements heavier than iron, including cherished ones like silver and gold, dangerous ones like radon and uranium, and ones you've never heard of (or could pronounce if you had) like praseodymium and ytterbium.

Two ways to you Stars have one of two ways to produce these heavier-than-iron elements—and, not incidentally, to get them and all the other elements forged in their nuclear furnaces out into space so they can be incorporated into new stars, planets, and people. Some of that widely dispersed stardust is holding you up right now. The first way occurs in red giants. These are stars that have burned up all the hydrogen in their centers. When that happens, the star becomes, as the astrophysicist Craig Wheeler has put it, somewhat schizophrenic: The core loses energy, contracts, and heats up even as the envelope—the rest of the star outside the core—gains energy, expands, and cools (and appears redder). The expansion is quite, well, expansive: When our sun becomes a red giant, it will grow so large that it will engulf and evaporate the inner planets, including the Earth. Some red giants last long enough to create elements in their cores heavier than iron through something called the s-process, for slow. Over a time scale of thousands of years, the s-process can result in the manufacture of elements all the way up to bismuth (83).

These get pulled to the star's surface by convection and sloughed off into space via the star's stellar wind. Some of that widely dispersed stardust is holding you up right now. This spectacular false-color image shows Cassiopeia A, the remnant of a supernova. At the center of the image lies the dead star, while surrounding it is the rapidly expanding shell of material blasted away from the star as it died. Enlarge Photo credit: NASA/JPL-Caltech/O. Krause (Steward Observatory) A real blast Elements heavier than bismuth only arise through the r-process, for rapid. How rapid? Seconds flat. The r-process is what happens when a star explodes in a supernova. It's easy for us to think of stars as lasting essentially forever, but the most massive stars survive only a few million years—a cosmic moment, really—and when they go, they go fast. What happens? When a red giant gets to the stage of having fused all its lighter elements and is left with an iron core, the star can no longer retain its equilibrium—heat energy pushing out as gravity pulls in. Gravity suddenly gains the upper hand, collapsing the core all at once to billions of times the density of the Earth.

The star then blows itself apart in an astronomical cataclysm. For a brief period, it shines as brightly as an entire galaxy and releases as much energy as our sun will in its 10-billion-year lifetime. In the first few seconds, protons in the atoms created during the star's life collide with highly energetic neutrons, fashioning in an instant all the naturally occurring elements heavier than bismuth up to uranium, and even a few short-lived still-heavier elements such as plutonium and californium. All these blast out into space at millions of miles an hour, seeding the interstellar medium with the atoms that eventually end up in new stars, new solar systems, and, in your case, you. In this view of the Carina Nebula, the Hubble Space Telescope captured a tumult of star birth and death. In the image, green corresponds to hydrogen, blue to oxygen, and red to sulfur—three of the 92 naturally occurring elements that space has bequeathed to us. Enlarge Photo credit: For Hubble Image: NASA, ESA, N. Smith (University of California, Berkeley), and The Hubble Heritage Team (STScI/AURA). For CTIO Image: N.

Smith (University of California, Berkeley) and NOAO/AURA/NSF The birth of you Over time, molecular clouds of gas and dust out in deep space develop from those strewn elements and begin to contract under their own gravity. Such clouds are almost all hydrogen and helium, but they've got a scatter of heavier elements, too. And the most abundant elements begin to assemble into molecules, simple ones like water (H2O) and more complex ones like the sugar glycoaldehyde (C2H4O2). Astronomers can identify these compounds, and individual elements, using spectrometers. Eventually, a kind of raw-clay star called a proto-star forms, with a disk of material surrounding it that will eventually beget planets. That process happened in our own solar system about five billion years ago, resulting in the sun, the planets, and, five billion years later, you. Did you have any idea you have atoms in your body that are over 13 billion years old? Just how those atoms and molecules that ended up on our planet went from non-living to living remains one of the great unanswered questions in science.

But where the elements came from to start with has now been worked out, in broad strokes anyway, to astrophysicists' widespread satisfaction. It is an amazing story, isn't it? Â *While still being sorted out, current estimates hold that, in time, stars below about eight solar masses collapse into white dwarfs, a very dense ex-star; those between about eight and 30 solar masses explode in supernovae (and release their heavy elements); and, with certain exceptions, those between about 30 and 100 solar masses collapse into black holes (and retain their heavy elements). Peter Tyson is editor in chief of NOVA Online. The author would like to thank Craig Wheeler, University of Texas at Austin, for help with this article. Sources

Question: Which is the lightest naturally occurring element?
Answer: HYDROGEN